copyright 2004 - john sayles1 bonding and geometry review unit 8 chapters 9 and 10

Copyright 2004 - John Sayles 1

Bonding and Geometry

Review Unit 8

Chapters 9 and 10


Chemical Bonding Atoms react in order to achieve a full outer shell

– “stability” refers to the inability to react

– Full outer shell is a low energy configuration

Three main strategies– Ionic bonding

• Metal and nonmetal, one gains/one loses, e-’s transferred

– Covalent bonding• Two nonmetals, both gain, e-’s shared

– Metallic bonding• Two metals, both lose, cations in sea of e-’s


Ionic Bonding Metal loses e-’s, forms pos. cations Nonmetal gains e-’s, forms neg. anion Ionic bond is the electrostatic attraction between the

cation and anion– Only one variation; no double, etc…– To compare the strength of ionic bonds, use Coulomb– Five steps of the Born-Haber cycle

• Form gaseous metal atoms, gaseous nonmetal atoms• Form cations, anions• Form ionic bonds and ionic crystal

Special valences– Trans metals lose s e-’s first (+2) then d’s (Sc,Fe,Ag,Zn)– Pseudo TM’s lose p’s first, then d’s (Pb+2 or +4)


Born-Haber Cycle for NaCI

Copyright © Houghton Mifflin Company. All rights reserved 9-3

9_3

Na(s)

Na(g)

+

+

Step 1

Step 2 Step 5

Cl2(g)

Cl(g) ClŠ(g)

NaCl(s)Direct

route

Step 4

Na+(g)Step 3

21


Properties of Ionic Compounds Crystalline, brittle

– Lattice energy = ionic bond = 400kJ/mol (Coulomb)

Solids at normal temps– VERY high MP’s and BP’s; no IMF’s, must break ionics

Soluble in polar solvents– “Ionic” is the ultimate in polar

– Ions are vigorously hydrated

Conduct in melt or solution– Chock full o’ ions

– Must be mobile, as in melt or sol’n


Covalent Bonding Both nonmetal atoms need to gain e-’s Impossible, so best option is to share e-’s

– Nuclei are mutually attracted to the shared pair Many variations

– Multiple bonds• Share 2 or 3 pairs of e-’s• Stronger and shorter• First bond is , others are

– Polar vs. nonpolar• ∆ electronegativity cause uneven sharing• Makes molecule polar if geometry is asymmetrical

– Resonance involves bonds flipping– Network and coordinate bonds


Covalent Bonding in the H2 Molecule


9_8


Covalent Bond Properties Bond strength depends on

– Order (# of shared pairs)– Polarity (∆ eneg)– Environment (eneg neighbors weaken bonds)– Indicated by bond energy (which, ironically, ignore

polarity and environment) Length “inversely proportional” to strength Vibrational frequency

– Stronger bonds vibrate at higher frequency, just like guitar strings

– Basis for spectroscopy


Bond Energies (in kJ/mol)

Copyright © Houghton Mifflin Company. All rights reserved Table 9.5

T9_5

Single Bonds

Multiple Bonds

C N O S

HCNOSFClBrI

H

432411386459363565428362295

346305358272485327285213

167201Ń283313ŃŃ

142Ń190218201201

226284255217Ń

155249249278

240216208

190175 149

F Cl Br I

CC CC CN NN N

C NC NN OO2

C OC OS O (in SO2)S O (in SO3)

602835418942

615887607494

745 (799 in CO2)1072

532469

*Data are taken from J. E. Huheey, Inorganic Chemistry, 4d ed. (New York: Harper & Row, 1993), pp. A21-A34


Properties of Covalent Cmpds Low MP’s and BP’s

– Only mess with IMF’s to melt or boil

Solubility depends on polarity of compound– Polar dissolve in polar solvents, vice-versa

Don’t conduct in any state, except for …– Acids


Lewis Structures Great way to represent the e- arrangement in a

covalent molecule– Do the Have/Need thing to count bonds

• Atoms “need” 8 e- each, except H

• “Have” refers to outer shell e-’s only

– (need - have) ÷ 2 = number of bonds

– Arrange atoms with highest bond capacity in middle

– Use sticks and snakebites to represent bonding and nonbonding e- pairs

Do NH3, CO2, N2, CO3-2, PF5


Metallic Bonding Both metals need to lose e-’s Both do lose, creating “sea of e-’s”

– Accounts for conductivity, flexibility, luster

Vary in strength, as do MP’s and BP’s Insoluble in all solvents

– Acids convert them to salts, which dissolve

Conduct in the solid and melt, with e- sea


Delocalized Bonding in Sodium Metal


9_15


VSEPR Valence Shell Electron Pair Repulsion model Great for determing geometry of molecule Doesn’t explain

– Equivalence of bonds made from non-equiv orbitals– Multiple bonding– Para and diamagnetism

Premises– Basic geometry due to need to spread e- pairs out– Nonbonding (lone) e- pairs need more room

• They’re delocalized

– Multiple bond treated as a single e- pair


Application of VSEPR Draw Lewis structure Total e- pairs determine basic geometry

– Linear, trigonal planar, tetrahedral, trigonal bipyramid, octahedral

Consider # and location of lone e- pairs to pick the specific geometry– Keep in mind lone pairs occupy roomiest

positions– See models


Basic Geometries



Specific Geometries

10-4ABCopyright © Houghton Mifflin Company. All rights reserved


More Specific Geometries

Copyright © Houghton Mifflin Company. All rights reserved 10-9AB


Valence Bond Model Supposes the formation of Hybrid Orbitals

– To explain the equivalence of bonds– Sp, sp2, sp3, dsp3, d2sp3

• Each hybridization corresponds to a basic geometry

– Hybrids hold 1st bonding e- pairs and lone e- pairs• All bonds start with the overlap of hybrids

– Forms a bond

• 2nd and 3rd bond due to overlap of unhybribized p’s– Forms a bond


Hybrid Orbitals and Their Spatial Arrangements



Bonding in H2O


10_25


Bonds vs. Bonds Bonds

– The overlap of hybrids– Overlap is along bond axis– Stronger (lower energy), approx. 400 kJ/mole– Can swivel

Bonds– Overlap of unhybridized p orbitals– Overlap is of bond axis (above and below)– Weaker, approx. 70% of bond (~280kJ/mol)– Cannot swivel

• Creates cis- and trans- isomerism (geometric isomerism)

– Involved in resonance• e-’s not localized


Sigma and Pi Bonds



Bonding in Ethylene



Bonding in Acetylene



Valence Bond Model Designed to explain equivalence of bonds made

from nonequivalent orbitals– Throw s’s and p’s in the quantum blender to form

equivalent hybrids Also explains details of multiple bonding

– No “wink-wink” bogosity– Explains relative strengths of single, double, triple bonds

bonds not as strong as bonds

– Explains resonance• e-’s in p orbitals are delocalized

Does NOT explain para and diamagnetism


Molecular Orbital Model Supposes the formation of molecular orbitals

– Not to be confused with hybrid atomic orbitals– Use QM to combine ’s atomic orbitals

Combining 2 atomic orbitals creates 2 molecular orbitals– One is lower in energy ---> bonding

• Putting e-’s here creates attraction

– One is higher in energy ---> anti-bonding• Putting e-’s here creates repulsion


Formation of Bonding and Anti-bonding Orbitals from 1s Orbitals to Hydrogen Atoms



Using the MO Model Create MO’s

– Draw the MO diagram (see following slides)

Fill MO’s using the Aufbau Principle– Lowest energy first

– Spread out e- among degenerate MO’s

– Only 2 e-’s per MO

Interpret MO diagram– Bond order = # bonding pairs - # antibonding pairs

– Paramagnetism requires at least one unpaired e-


Relative Energies of 1s of the H atom and the sigma1s and sigma*1s Molecular Orbitals of H2


10_32

Ene

rgy 1s

1s

1s

1s

H atom H atomH2 molecule


The Energies of the Molecular Orbitals of Li2


10_33

Li atom Li atom

En

erg

y2s

2s

2s

2s

1s

1s

1s

1s

Li2 molecule


Relative Energies of Molecular Orbitals of Homonuclear Diatomic Molecules (Excluding K Shells)


10_35

2p 2p

2p

2p

2s

*2p

*2p

2s 2s*2s

Atom AtomMolecule

Ene

rgy


Molecular Polarity In order for a molecule to be polar, it must

– Contain polar bonds• ∆ electronegativity

• Most bonds are at least a little polar

• Measured as dipole moment (in DeByes’s)

– And be asymmetrical

Polarity affects properties– IMF’s

– MP, BP, volatility, solubility, viscosity, surface tension

copyright 2004 - john sayles1 bonding and geometry review unit 8 chapters 9 and 10

Documents